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Chem 102 Chapter 12

by: Saida Muktar

Chem 102 Chapter 12 CHEM 102 001

Saida Muktar

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Chapter 12 notes
Fundamental Chemistry II
Dr. Perks
Class Notes
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This 11 page Class Notes was uploaded by Saida Muktar on Tuesday February 16, 2016. The Class Notes belongs to CHEM 102 001 at University of Maryland - Baltimore County taught by Dr. Perks in Spring 2016. Since its upload, it has received 35 views. For similar materials see Fundamental Chemistry II in Chemistry at University of Maryland - Baltimore County.

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Date Created: 02/16/16
12.1 The Condensed Phases  At a given temperature, the magnitude of intermolecular forces determines whether a substance is a solid, a liquid, or a gas. 12.2 Properties of Liquids  Several of the physical properties of a liquid depend on the magnitude of its intermolecular forces. 1 Surface Tension o A molecule within a liquid is pulled in all directions by the intermolecular forces between it and the other molecules that surround it. There is no net pull in any one direction. o A quantitative measure of the elastic force in the surface of a liquid is the SURFACE TENSION, the amount of energy required to stretch or increase the surface of a liquid by a unit area. o A liquid with strong intermolecular forces has a high surface tension. o A meniscus is another illustration of surface tension which is the curved surface of a liquid contained in a narrow tube. The surface tension of water causes this film to contract, and pulls the water up the cylinder, this effect is known as CAPILLARY ACTION.  Two types of forces bring about capillary action.  COHESION, the attraction between like molecules  ADHESION, the attraction between unlike molecules.  If adhesion is stronger than cohesion then the contents of the tube will be pulled upward. 2 Viscosity o 2 VISOCITY, with units of N.s/m , is a measure of a fluid's resistance to flow. o The higher the viscosity, the more slowly a liquid flows. o The viscosity of a liquid typically decreases with increasing temperature. o Liquids that have strong intermolecular forces have higher viscosities than those that have weaker intermolecular forces. 3 Vapor Pressure of Liquids o Vapor Pressure is another property of liquids that depends on the magnitude of intermolecular forces. o The vapor pressure over the liquid increases until the rate of condensation is equal to the rate of evaporation, which is constant at any given temperature.  This situation, wherein a forward process and reverse process are occurring at the same rate, is called a DYNAMIC EQUILIBRIUM.  Although both processes are ongoing, the number of molecules in the gas phase at any given point in time does not change.  The pressure exerted by the molecules that have escaped to the gas phase, once the pressure has stopped increasing, is the EQUILIBRIUM VAPOR PRESSURE. o At a higher temperature a greater percentage of molecules at the liquid surface will possess sufficient kinetic energy to escape into the gas phase. o A linear relationship exists between the natural log of vapor pressure and the reciprocal of absolute temprature.  This relationship is called the Clausius-Clapeyron equation: 12.3 The Properties of Solids  The magnitudes of intermolecular forces in solids are responsible for some of their important physical properties, including melting point and vapor pressure.  Solids fall generally into two categories: amorphous, and crystalline 1 Melting Point o The MELTING POINT of a solid is the temperature at which the energies of individual particles enable them to break free of their fixed positions in the solid -- allowing them to flow past one another. o At the melting point, the solid and liquid phases of a substance coexist in equilibrium. o Solid substances with strong intermolecular forces melt at higher temperatures than those with weaker intermolecular forces. o Attractive forces are greater between larger molecules than smaller molecules; greater between polar molecules than between nonpolar molecules; and greater between molecules that form hydrogen bonds than between molecules of similar size that cannot for hydrogen bonds.  Because ion-ion interactions are typically much stronger than other types of intermolecular forces, the melting points of ionic solids tend to be very high. 2 Vapor Pressure of Solids o Solids have characteristic vapor pressures that depend on the magnitude of intermolecular forces.  Because molecules are more tightly held in a solids, the vapor pressure of a solid is generally much lower than that of the corresponding liquid. o The vapor pressures of solids are typically very low at room temperature. 3 Amorphous Solids o In most solids, the atoms, molecules, or ions occupy positions in a regular three-dimensional arrangement.  These substances are referred to as CRYSTALLINE; and they constitute the majority of solids because of their inherent stability.  If a solid forms under certain extraordinary conditions, there may not be sufficient time for the atoms or molecules to move into the positions of a regular crystal before they become locked in place. The resulting substances are known as AMORPHOUS SOLIDS. 4 Crystalline Solids o A CRYSTALLINE SOLID possesses rigid and long-range order, its atoms, molecules, or ions occupy specific positions.  The arrangement of particles in a crystalline solid, which we call the LATTICE STRUCTURE, depends on the nature and the size of the particles involved. o The forces responsible for the stability of a crystal can be ionic forces, covalent bonds, van der Waals forces, hydrogen bonds, or a combination of some of these forces. o A UNIT CELL is the basic repeating structural unit of a crystalline solid. 1/29/2016 8:28 PM - Screen Clipping o Every crystalline solid can be described in terms of one of the seven types of unit cells. 1/29/2016 8:30 PM - Screen Clipping  The geometry of the cubic unit cell is particularly simple because all sides and all angles are equal.  Any of the unit cells, when repeated in space in all three dimensions, forms the lattice structure characteristic of a crystalline solid. o Packing Spheres  We can understand the geometric requirements for crystal formation by considering the different ways of packing a number of identical atoms to form an ordered three-dimensional structure.  The way the atoms are arranged in layers determines the type of unit cell.  The three-dimensional structure can be generated by placing a layer above and below this layer in such a way that atoms in one layer are directly over the atoms in the layer below. 1/29/2016 8:46 PM - Screen Clipping  The COORDINATION NUMBER is defined as the number of atoms surrounding an atom in a crystal lattice.  The value of the coordination number indicates how tightly the atoms are packed together -- the larger the coordination number, the closer the atoms are to one another.  The basic repeating unit in the array of atoms is called a SIMPLE CUBIC CELL.  The other types of cubic cells are the BODY-CENTERED CUBIC CELL and the FACE-CENTERED CUBIC CELL. 1/29/2016 8:50 PM - Screen Clipping 1/29/2016 10:07 PM - Screen Clipping 12.4 1/29/2016 10:29 PM - Screen Clipping Types of Crystalline Solids  The structures and properties of crystalline solids, such as melting point, density, and hardness, are determined by the kinds of forces that hold the particles together. o We can classify any crystal as one of the four types: ionic, covalent, molecular, or metallic 1 Ionic Crystals o Ionic crystals are composed of charged spheres that are held together by Coulombic attraction.  Anions typically are considerably bigger than cations and the relative sizes and relative numbers of the ions in a compound determine how the ions are arranged in the solid lattice. o Most ionic crystals have high melting points, which is an indication of the strong cohesive forces holding the ions together. o A measure of the stability of ionic crystals is the lattice energy; the higher the lattice energy, the more stable the compound. o Ionic solids do not conduct electricity well because the ions are fixed in position. 2 Covalent Crystals o In covalent crystals, atoms are held together in an extensive three- dimensional network entirely by covalent bonds. 3 Molecular Crystals o In molecular crystals, the lattice points are occupied by molecules, so the attractive forces between them are can der Waals forces and/or hydrogen bonding. 4 Metallic Crystals o Every lattice point in a metallic crystal is occupied by an atom of the same metal. o Metallic crystals are generally body-centered cubic, face-centered cubic, or hexagonal close-packed. o Metallic elements are usually very dense. 12.5 Phase Changes  A phase is a homogeneous part of a system that is separated from the rest of the system by a well-defined boundary. o When a substance goes from one phase to another phase, we say that it has undergone a PHASE CHANGE.  Phase changes in a system are generally caused by the addition or removal of energy, usually in the form of heat. 1/29/2016 10:12 PM - Screen Clipping 1 Liquid-Vapor o Because boiling point is defined in terms of the vapor pressure of the liquid, the boiling point is related to the MOLAR HEAT OF VAPORIZATION( Δh vap) the amount of heat required to vaporize a mole of substance at its boiling point.  The boiling point generally increases as Δh vapincreases.  Bothe the boiling point and Δh vapare determined by the strength of intermolecular forces. o The opposite of vaporization is condensation.  A gas can be liquefied (made to condense) either by cooling or by applying pressure.  Cooling a sample of gas decreases the kinetic energy of its molecules, so ev3entually the molecules aggregate to form small drops of liquid.  Applying pressure to the gas (compression) reduces the distance between molecules, so they can be pulled together by intermolecular attractions.  Many liquefaction processes use a combination of reduced temperature and increased pressure. o Every substance has a CRITICAL TEMPRATURE (T ) above whCch its gas phase cannot be liquefied, no matter how great the applied pressure.  T is the highest temperature at which a substance can exist as c a liquid. o CRITICAL PRESSURE (P ) is tCe minimum pressure that must be applied to liquefy a substance at its critical temperature. o At temperatures above the critical temperature, there is no fundamental distinction between a liquid and a gas - we simply have a fluid.  A fluid at a temperature and pressure that exceed critical temperature and critical pressure is called a SUPERRITICAL FLUID. 1 Solid-Liquid o The transformation of a solid to a liquid is called melting, or FUSION.  For a given substance, the melting point and freezing point are the temperature at which solid and liquid phases coexist in equilibrium.  The melting point and freezing point at 1 atm specifically are called the normal melting point and normal freezing point. 1/29/2016 10:16 PM - Screen Clipping o The MOLAR HEAT OF FUSION (ΔHfusis the energy, usually expressed in kJ/mol, required to melt 1 mole of a solid o If we remove heat from a gas sample at a steady rate, its temperature decreases.  As the liquid is being formed, heat is given off by the system, because its potential energy is decreasing. o SUPERCOOLING is a phenomenon in which a liquid can be temporarily cooled to below its freezing point. Super cooling occurs when heat is removed from a liquid so rapidly that the molecules literally have no time to assume the ordered structure of a solid. A super cooled liquid is unstable. 1 Solid-Vapor o Solids can be vaporized and so have a vapor pressure. o SUBLIMATION is the process by which molecules go directly from the solid phase is the vapor phase.  The reverse process, in which molecules go directly form the vapor phase to the solid phase is called DEPOSITION. o Because molecules are more tightly held in a solid, the vapor pressure of a solid is generally much less than that of the corresponding liquid. o The MOLAR ENTHALPY OF SUBLIMATION (ΔHsub of a substance is the energy required to sublime 1 mole of a solid.  It is equal to the sum of the molar enthalpies of fusion and vaporization: 12.6 Phase Diagrams o A PHASE DIAGRAM summarizes the conditions (temperature and pressure) at which a substance exists as a solid, liquid or gas. The graph is divided into three regions, each of which represents a pure phase The line separating any two regions, called a phase boundary line, indicates conditions under which these two phases can exist in equilibrium. The point at which all three phase boundary lines meet is called the TRIPLE POINT.  The triple point is the only combination of temperature and pressure at which all three phases of a substance can be in equilibrium with one another. The point at which the liquid-vapor phase boundary line abruptly ends it’s the critical point, corresponding to the critical temperaturec(T ) and the critical pressure (c ).


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